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CE610 ANALYSIS OF NATURAL AND

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					 CE610 ANALYSIS OF NATURAL AND
        POLLUTED WATERS

              LABORATORY MANUAL


                 Autumn Quarter, 2006




Civil and Environmental Engineering and Geodetic Science
                The Ohio State University
                 Columbus, OH 43210
                                            TABLE OF CONTENTS



Periodic Table of the Elements …………………………………………………….………1

Laboratory Notebook and Grading.................................................................................... 2

Laboratory Safety and Standard Practice ........................................................................... 3

City of Columbus Water Quality Data …………………………………………………….8

Laboratory Experiments

         Lab. #1 Laboratory Safety and Quality Assurance /Quality Control (QA/QC) ....... 9

         Lab. #2 Solids Characteristics of Water .............................................................. 11

         Lab. #3 Standardization of Acids and Bases ........................................................ 18

         Lab. #4 Determination of pH and Alkalinity ....................................................... 22

         Lab. #5 Titrimetric Determination of Hardness ................................................... 27

         Lab. #6 Colorimetry of Phosphorus..................................................................... 31

         Lab. #7 Biological Oxygen Demand ................................................................... 36

         Lab. #8 Colorimetry of Iron ................................................................................ 41

 Acknowledgment .......................................................................................................... 44
                                                            Periodic Table of the Elements


 1.00794                                                                                                                                                                                       4.002602

  H                                                                                                                                                                                             He
   1                                                                                                                                                                                              2
  6.941     9.012182                                                                                                                   10.811     12.011    14.00674    15.9994    18.998403 20.1797

  Li         Be                                                                                                                         B          C          N           O           F         Ne
   3           4                                                                                                                         5          6          7           8          9          10
22.989768   24.3050                                                                                                                   26.981539 28.0855     30.973762    32.066     35.4527     39.948

 Na         Mg                                                                                                                          Al         Si         P           S          Cl         Ar
   11         12                                                                                                                        13          14         15         16          17         18
 39.0983     40.078    44.955910    47.88     50.9961     51.9961    54.93805   55.847   58.93320     58.69     63.546      65.39      69.723      72.61    74.92159     78.96      79.904      83.80

   K         Ca         Sc           Ti         V          Cr        Mn         Fe Co                 Ni        Cu          Zn         Ga Ge                 As          Se          Br         Kr
   19         20          21         22         23          24         25        26         27         28         29         30         31          32         33         34          35         36
 85.4678     87.62     88.90585     91.224    92.90638     95.94     98.9063    101.07   102.90550   106.42    107.8682    112.411     114.82     118.710    121.75      127.60    126.90447    131.29

 Rb          Sr          Y          Zr         Nb Mo                  Tc        Ru Rh                Pd         Ag         Cd           In        Sn         Sb          Te           I         Xe
   37         38          39         40         41          42         43        44         45         46         47         48         49          50         51         52          53         54
132.90543   137.327                 178.49    180.9479     183.85    186.207    190.2     192.22     195.08    196.96654   200.59     204.3833     207.2    208.98037 208.9824 126.90447 222.0176

 Cs          Ba        La-Lu        Hf         Ta          W          Re        Os          Ir        Pt        Au         Hg           Tl        Pb          Bi         Po          At         Rn
   55         56        57-71        72         73          74         75        76         77         78         79         80         81          82         83         84          85         86
223.0197    226.0254               261.1087   262.1138    263.1182   262.1229

  Fr         Ra        Ac-Lr Rf/Ku Ha/Ns                  Unh Uns Uno Une
                       89-103       104         105
   87         88                                           106        107       108        109


                       138.9055    140.115    140.90765    144.24    146.9151   150.36    151.965     157.25   158.92534    162.50    164.93032   167.26    168.93421    173.04     174.967

                         La         Ce          Pr         Nd Pm Sm Eu                               Gd Tb                  Dy         Ho          Er        Tm Yb                   Lu
                          57         58          59         60         61        62         63         64         65         66          67         68         69         70          71
                       227.0278    232.0381   231.0359    238.0289   237.0482 244.0642   243.0614    247.0703 247.0703     251.0796   252.0829 257.0951     258.0986    258.0986   260.1053

                         Ac         Th         Pa           U         Np Pu Am Cm Bk                                        Cf         Es Fm Md No                                   Lr
                          89         90          91         92         93        94         95         96         97         98          99        100        101        102         103




                                                                                                                                                                                                          1
                           LABORATORY NOTEBOOK AND GRADING



The laboratory notebook is used to keep a permanent record and diary of your laboratory work. It should
be regarded as a legal document, like survey notebooks, because it is the basis of patent claims, legally
required reports to the government, client billings and timesheets, and is the principal means of retracing
work for errors and checking.

Notebook Requirements: Use a notebook with the following characteristics.

           pages glued and stitched into the binding (do not use loose-leaf notebooks or loose
            papers of any kind)

           duplicate removable pages and carbon sheets, allowing the student to make copies
            directly when data and notes are written down (laboratory notebooks utilizing
            carbonless duplicate pages are also acceptable)

Preparation of Laboratory Reports. Each week you must submit a laboratory report completed in your
laboratory notebook. A report is required for every lab during the quarter and must include the following
sections: Title, Purpose, Materials and Methods, Pre-Lab Questions, Results, and Post-Lab Questions.
The title, purpose, and materials and methods section, as well as all pre-lab questions, should be
completed prior to the start of class as part of your “pre-lab”. All laboratory work must be completed in
the laboratory notebook and loose-leaf papers are not acceptable.

Pre-Lab Quiz. Each laboratory period starts with a quiz that covers the pre-lab.

Recording Data. RECORD EVERYTHING! All measurements and arithmetic must be directly
written in the laboratory notebook. Use a non-erasable, oil-based ink ballpoint pen and cross out errors
with a single stroke that still enables the error to be read. Never erase or use “Whiteout” to cover up
errors. This is a legal requirement in many studies (e.g., patents), simplifies error checking, and
eliminates the suspicion that the laboratory work was faked. Oil-based inks do not smear or run when
chemicals are spilled on them and the recorded work is not lost unless the chemical destroys the paper
itself.

Date, title, and sign (or initial) each page of the notebook. This is a legal requirement for patent applica-
tions and for data introduced as evidence in court cases. Some laboratories use “bench sheets” that are
special forms for data entry. In this case, all data and arithmetic are done directly on the bench sheets or
in specifically designated notebooks. All bench sheets are kept permanently in binders. In addition,
electronic laboratory notebooks are used in some situations. The main disadvantage with using electronic
notebooks, however, is the inability to maintain an unalterable, permanent record of the work done.

Submit. Turn in the yellow (or blue) carbon copy sheets from your laboratory notebook for each week’s
lab. Each laboratory report must be turned in at the beginning of the following laboratory meeting to
receive full credit. Reports not turned in at this time will receive an automatic 10 percent deduction. For
each subsequent day the assignment is late (weekends included) the score will be reduced by an
additional 10 percent.

Grading. The laboratory reports and quizzes are worth 20% of your overall class grade (15% for the
laboratory reports and 5% for the pre-lab quizzes).




                                                                                                           2
                          LABORATORY SAFETY AND STANDARD PRACTICE

An analytical chemistry laboratory can be a dangerous place if one is careless. The frequency of acci-
dents and their severity can be greatly reduced if one uses common sense and caution and follows the
rules below.

                                     Emergency Telephone Number
                                      For fire and personal injury
                                          dial: 911 or 2-2525

You are in room 008, Hitchcock Hall, 2070 Neil Ave. The telephone is located in this room next to the
lab entrance. There are three exists from the building; two are accessible from the south end of the hall
and one can be reached from the north end.

Fire Extinguishers. There are three (3) extinguishers located within close proximity to the laboratory:
                south end of the hallway, just to the right of the lab entrance
                north end of the hallway, opposite room 050
                just north of 008, next to room 025

The fire extinguishers in the laboratory are intended for use on trash, wood, and paper fires (Class A),
grease and oil fires (Class B), and electrical fires (Class C). In the case of small fires, remove the fire
extinguisher from the wall bracket and pull out the safety pin in the handle. Point the nozzle at the base
of the fire and squeeze the handle.

Fire Alarms. In the case of large fires, activate the fire alarm and leave the building. Fire alarms are
located at the south end of the hallway next to the exit doors and at the north end of the hallway, opposite
room 051.

Eyewash Fountains and Safety Shower. There is an eyewash fountain and safety shower located near
the center of the laboratory. The eyewash consists of two fixed spray heads attached to either side of a
small round sink or two valves on the sides of the sinks. To operate, bend over the sink and turn on the
valve on the same side as the affected eye.

If you are splashed with a corrosive chemical, go to the emergency shower, stand under it, and pull the
chain. You will be drenched with water, which will remove and dilute the chemicals in your clothing and
on your skin and hair. Acids are absorbed and retained by cloth and leather. If acid is spilled on clothing
or shoes, remove and change them as soon as possible and have them thoroughly cleaned before reuse.

If chemicals splash or spray into your face or eyes, immediately remove the safety goggles and any other
eyeglasses you may be wearing and spray the affected eyes. Hold the lids open and continue spraying for
at least 15 minutes. Contacts are expressly forbidden!

Note: Alkali burns are much more serious to the eye than acid burns because alkali penetrates the eye. In
case of alkali accidents, irrigation of the eye must begin immediately, within 10 to 30 seconds.

First Aid Kits. Located near the entrance to room 008.

Protective Clothing. Safety goggles and aprons must be worn at all times while you are in the
laboratory. Goggles, aprons, and thermally insulated gloves must be worn when handling hot materials,
e.g. placing or removing dishes in or from the drying ovens and muffle furnaces. Goggles, aprons, and
rubber gloves must be worn when cleaning glassware.


                                                                                                          3
Do not wear contact lenses in the laboratory. Gases and vapors accumulate under the lenses and can
damage the cornea. Soft lenses absorb solvent vapors, even if protective goggles are worn. Also, contact
lenses reduce the effectiveness of eye washing.

Wear old clothing. If you wouldn’t paint a house or play rugby in it, don’t wear it! Spills are certain, and
most of the chemicals in the laboratory corrode and/or stain clothing and leather. Wear leather shoes, not
sandals or cloth sneakers. Spilled chemicals can easily penetrate fabric shoes and socks, but not leather.
Leather shoes also provide protection against broken glass on the floor.

Chemical Data. Environmental Health and Safety (EHS) at OSU maintains a web page that allows
access to Material Safety Data Sheets (MSDS). An MSDS provides important information on properties
and appropriate controls for a chemical of interest. All known chemicals have an MSDS.

Spills. At some time, a container of acid or base may be dropped or turned over, and the contents may
spill out. If this happens you should first alert your fellow lab members and upon identifying the spilled
chemical follow the appropriate clean-up procedures.

Acid spills. Apply sodium bicarbonate solution to the entire spill to neutralize it and then wipe it up with
an absorbent material like a paper towel. Wear rubber gloves to avoid chemical burns and glass cuts and
dispose of the absorbent material in the indicated container. If there is any broken glass, dispose of it in
the broken glass container. Iron and aluminum solutions are acidic, and spills of these substances should
be treated as acid spills.

Base spills. Apply a boric acid solution to the entire spill to neutralize it and then wipe it up with an
absorbent material like a paper towel. Wear rubber gloves to avoid chemical burns and glass cuts and
dispose of the absorbent material in the indicated container. If there is any broken glass, dispose of it in
the broken glass container.

If a thermometer is broken, there will be a mercury spill. Mercury is a cumulative poison. It is volatile
and the vapors can be inhaled almost immediately upon exposure to the air. It can also be absorbed
through the skin. Do not touch it. Carefully observe where the mercury goes and inform the laboratory
instructor of the spill immediately. The instructor will use a special mercury clean-up kit.

Dilution of Acids and Bases. Many chemicals react violently when mixed. This is especially true of
concentrated acids and water. ALWAYS ADD ACID TO WATER; NEVER ADD WATER TO ACID!

Handling Glassware. Laboratory glassware is fragile and once broken can cause very serious wounds
that require immediate medical attention. Broken glassware may also be coated with dangerous
chemicals. Furthermore, the contents of glassware may be spilled.

Lubrication. Lubricate glass tubing before inserting into stoppers or plastic tubing. Suitable lubricants
include deionized water, glycerol, or a soap solution. Wrap the glass tubing in a folded paper towel and
grasp the tubing close to the stopper or plastic tubing. Insert using a rotating motion with very little
pressure. Push along the tubing axis, do not bend the tubing! If the force required seems excessive then
the glass tubing is too big. Thus, either select a smaller size of glass tubing or increase the diameter of
hole into which you are inserting the tubing. Thermometers and pipettes should be treated like glass
tubing.

Pipettes. Do not suck on pipettes. You may ingest some poisonous chemicals or your saliva may
contaminate the solution. Use the suction bulb (propipette) as instructed. Treat pipettes like glass tubing
and be sure to lubricate the top of the pipette before inserting into the suction bulb. Pipettes are fragile



                                                                                                          4
with respect to bending. Do not grip the pipettes between the fingers and the palm of the hand with the
thumb over the top. This causes bending. Instead, hold the pipette between the thumb and the second
and third fingers, and use the index finger to close the top. Never heat a pipette, which may change its
calibration and weaken the glass. Hold glass tubing and thermometers the same way.

Burettes. Improperly filling a burette may lead to spills down the burette sides or down your arm. The
spilled solution may damage counter tops or clothing and cause skin irritation or burns. To fill a burette,
first locate the burette stand on a low table or in a sink so that you may easily reach the top of the burette.
Place a small beaker under the burette to collect spills and verify that the stopcock is closed. Insert a
funnel into the top of the burette and slowly fill the burette to a point somewhat above the zero mark.
Carefully drain the burette down to or slightly past the zero mark.

Beakers. Beakers about 500 mL or smaller may be held in one hand wrapped around the sides (do not
squeeze), but larger breakers should be held with two hands, one underneath and one around the sides.
Beakers are usually broken when they are set down hard on a laboratory bench or on top of other equip-
ment. Set them down softly and make sure the area is clear. Hot beakers should be handled with ther-
mally insulated gloves and tongs and set down on ceramic pads.

Volumetric flasks. Volumetric flasks have long, fragile necks that are easily broken when used as
handles for shaking. Use two hands, placing one on the bulb and the other on the neck, when shaking a
flask to mix its contents.

Do not use volumetric flasks to store chemicals. They are fragile and their narrow bottoms make them
unstable and liable to tip over. Never heat volumetric glassware, which may alter the calibration and
weaken the glass.

Heating and cooling glassware. Laboratory glassware, even Pyrex, may shatter if it is rapidly or unevenly
cooled after heating. Starred, chipped, cracked or scratched glassware is especially liable to thermal
breakage. Even if the glassware does not break on cooling, it may retain residual stresses and be weaker
than before.

Do not heat soft glass, especially thick-walled apparatus, because soft glass easily deforms and is very
susceptible to thermal stock. Soft glass is normally found in solution storage bottles, volumetric
glassware (pipettes and flasks), funnels, jars, droppers, watch glasses, desiccators, glass plates and test
tubes. Hot glass cools very slowly. Use tongs, grips or thermally insulated gloves when handling heated
glassware. Hot alkali dissolves laboratory glassware. Do not heat or boil alkali solutions.

Broken glassware. Broken glassware should be disposed in designated containers. The janitors do not
expect to find broken glass in normal trash containers and they may be injured. Also, only broken glass
should be placed in the broken glass containers.

Cleaning. Clean your laboratory space and equipment at the end of each laboratory session. Contami-
nated laboratory glassware is very difficult to clean once it has dried and the contaminants have set. Also,
the space you are using is shared by others, who do not know what you were doing. Do not leave
surprises behind.

Cleaning glassware. Goggles, aprons, and rubber gloves must be worn when cleaning glassware. This
protects against both chemical burns and lacerations due to glass breakage.

The cleanliness of glassware is judged by the absence of visible spots and strains and by how water drains
from it. Hold the glassware upside down and watch how the water drains out. If the glass surface is



                                                                                                             5
clean, the water film gradually becomes thinner and the film is distributed uniformLy over the glass. If
the water film beads or streaks, the surface is not clean.

For most uses, glassware is cleaned by soaking in a warm 1 to 2% solution of special laboratory deter-
gent. Difficult deposits are removed by brushing. Do not use abrasive cleaning agents or tools because
they scratch the glass, which weakens the glass and leads to crevices which accumulate contaminants.
After washing, rinse glassware 5 times in tap water, 3 times in deionized water, and set to drain and air
dry. The tap water rinse may be applied as a spray from a hose. If tap water does not remove the entire
detergent residue, rinse the glassware two or three times in dilute hydrochloric acid followed by three
rinses in deionized water. Note that some detergents contain trisodium phosphate, which is a very strong
alkali and can cause serious skin burns and eye damage.

In case of very stubborn dirt, an acid wash can be used. The wash solution is prepared by dissolving 10 g
of technical-grade sodium dichromate (Na2Cr2O7) in 200mL of hot, concentrated sulfuric acid (H2SO4).
When the sodium dichromate crystals are dissolved, the solution is cooled and placed into 500mL
decanters. Alternately, prepare a saturated aqueous solution of sodium dichromate (some crystals should
be left) and add 35mL of the saturated solution slowly and carefully (to avoid splattering) to a 9 lb bottle
of concentrated sulfuric acid. The sulfuric acid/dichromate solution is initially red-brown and becomes
greenish as the oxidant is consumed. This provides a convenient endpoint indicator.

Glassware to be acid cleaned is first rinsed in deionized water, drained and air-dried in order to (a)
prevent violent reactions between contaminants and the acid wash, and (b) prevent dilution of the acid
wash. After rinsing, pour the acid into the apparatus, coating all surfaces, and allow it to stand for several
minutes. When clean, invert the glassware over a plastic dishpan and drain off the acid wash for at least
20 minutes. If the solution is still reddish in color, the acid wash is later returned to its original container
for reuse. If it is greenish in color, it is discarded. Rinse the drained glassware 10 times with tap water,
then 3 times with deionized water, and finally set it out to drain and air-dry. To clean small pieces of
glassware submerge them in the acid wash.

Be very careful. The sodium dichromate/sulfuric acid wash is quite hazardous and will rapidly cause
second/third degree burns on the skin. If you come in contact with it, immediately rinse exposed skin
with large amounts of cool water. Also, both hexavalent and trivalent chromium ions are toxic to wildlife
and their discharge to sewers is restricted. Chromium can be recovered by reducing with sodium
thiosulfate, precipitation with sodium carbonate and filtration (see Coyne, G.S. 1992 The Laboratory
Handbook of Materials, Equipment and Technique, Prentice Hall, Inc., Englewood Cliffs, p. 209). If the
laboratory does not have a chromium recovery system, an alternative non-metallic oxidizer like
NOCHROMIXTM (Godax laboratories, NY, NY 10013) should be used.

Cleaning polyethylene. The procedure for cleaning polyethylene is somewhat different than that for
cleaning glass. Again, the usual procedure is to wash the polyethylene in a warm solution of a strong
laboratory detergent (using a brush if necessary), rinse 5 times in tap water, 3 times in deionized water,
and set it out to drain and air-dry.

If stubborn dirt remains, an acid wash is needed. However, the sodium dichromate/sulfuric acid wash
described above oxidizes polyethylene. Instead, use a solution of concentrated hydrochloric acid, and
perform the cleaning in a fume hood. Fill or submerge the polyethylene in the acid and let it soak for
several minutes. Drain the acid from the polyethylene back into its storage bottle, rinse the polyethylene
5 times in tap water, 3 times in deionized water, and set it out to drain and air-dry.

Waste Chemicals. Almost all the chemicals used in this course are acids or bases and can be disposed of
by dilution. Use the designated waste disposal sink that is located next to the water purification system.


                                                                                                              6
Run the cold water tap in the designated waste disposal sink, and gradually place small quantities of the
waste material into the flow near the drain hole. Be careful not to splash the chemicals. Concentrated
chemicals and solids may liberate enough heat to boil, spatter, or even cause steam explosions if too much
is added to the water. Concentrated solutions that fume (e.g., hydrochloric acid and ammonia) should be
disposed of in a fume hood sink. The dilute standard solutions and the titrated solutions may be disposed
of in your bench sink.

Food and Drink. Eating, drinking, and smoking in the laboratory are strictly forbidden because the air,
surface of benches, tables, chairs, and your hands and clothing may be contaminated with many
chemicals, some of which are poisonous. After each laboratory session, go directly to the nearest
appropriate restroom and wash your hands. Do this before eating, drinking, or smoking

Good Practice. The following is a few general guidelines to help ensure a safe and productive laboratory
environment.

   Buddy system. Nobody may work alone in the laboratory. In case of an accident, you will need
    someone to provide first aid and call for assistance.
   Use of chemicals. Dispense small quantities of chemicals from the original container. You can
    always take more if needed. Use a stainless steel spatula to remove chemicals from containers and
    place the chemicals in clean weighing papers, aluminum weighing dishes, mortars, etc. Never return
    unused chemicals to the original container. Strict adherence to this rule will prevent contamination of
    the laboratories chemical supplies and the production of spurious data. Dispose of unused chemicals
    safely, and in accordance with laboratory rules.
   Labeling. All containers must be labeled with

                1.      Name(s)
                2.      Date of preparation
                3.      Contents (if it is a solution, be sure to include the concentration).




                                                                                                         7
                                                  City of Columbus Water Quality Data

                                                Scioto River1                    Dublin Road WTP2                   Hap Cremean WTP2
Parameter                               (average)         (range)          (average)         (range)           (average)       (range)
Turbidity (NTU)                           85.29        15.84-218.55           0.35          0.05-0.35             0.12        0.09-0.21
Conductivity (S/cm)                       542            380-740             510            390-700              288          240-380
Total Dissolved Solids (mg/L)              378            200-460             318            170-410              210          160-290
Total Alkalinity (mg/L as CaCO3)           143            123-169              48             38-64                30            24-38
Total Hardness (mg/L as CaCO3)             242            174-280             120            116-124              105           89-125
Total Phosphate (mg-P/L)                   0.21          0.05-0.60            0.27          0.02-0.43             0.26        0.15-0.36
Nitrate (mg/L)                             4.05           0.7-6.4             4.09           1.2-6.2               1.5          0.7-2.6
pH                                          8.0           7.8-8.1              7.7           7.7-7.8               7.8          7.7-7.9
Total Organic Carbon (mg-C/L)              7.35         5.21-13.02            2.26          1.91-2.79             2.85        2.50-3.33
Fe (g/L)                                 1895           233-4824              50           <50-140               <50          <50-123

1
Inlet to the Dublin Road WTP (City of Columbus, Department of Public Utilities, Division of Water, Operations Report, 1996).
Compiled from 1996 – 2003 discharge data.
2




                                                                                                                                          8
                                          LABORATORY #1

                               LABORATORY SAFETY AND
                     QUALITY ASSURANCE /QUALITY CONTROL (QA/QC)

Purpose

The goals of this laboratory are to (1) produce a floor plan of the laboratory area indicating all relevant
safety features, and (2) develop an understanding of quality assurance and quality control.

Background

When performing chemical analyses of natural waters, the goal is to obtain accurate and reliable data in a
safe and efficient manner. Guidelines for performing chemical tests safely are provided at the beginning
of this manual. Below, important aspects of QA/QC are presented.

Quality Assurance. Quality assurance is defined as a set of operating principles that provide reliable data
of known quality. These operating principles are spelled out in a QA Plan. A QA Plan typically includes
the following: cover sheet with plan approval signatures, staff organization and responsibilities, sample
control and documentation, standard operating procedures for all analytical methods, calibration
procedures, quality control activities, performance audits, data assessment procedures, data reduction,
validation and reporting.

Quality Control. A good quality control program consists of the following elements: certification of
operating competence, recovery of known additions (“spikes”), analysis of externally supplied standards,
analysis of reagent blanks, calibration of standards, analysis of duplicates, and maintenance of control
charts. Quality control activities are integrated in all subsequent laboratory procedures to provide you the
opportunity to become familiar with these concepts. Descriptions of QC activities are provided below.

Recovery of known additions, or spikes, is used to evaluate the importance of matrix effects. The matrix
is the background liquid, solid, or gas in which the chemical of interest is present, and compounds present
in the matrix may interfere with the analysis procedure. For example, the presence of high levels of
dissolved solids can interfere with trace element analysis by inductively coupled plasma mass
spectrometry (ICP-MS). The precision for known additions is expressed as the percentage of the known
addition recovered. The expected recovery of the known addition depends on the chemical of interest, but
is usually in the range of 80%-120%. At least 10% of samples analyzed should be spikes.

The analysis of externally supplied standards is also an important QC activity. The use of external
standards provides verification for internal calibration procedures. Internal calibration standards can be
incorrectly prepared, and can subsequently influence the accuracy of the method. Thus, analyzing
externally supplied samples provides a measure of the precision of the method. Both NIST and USEPA
provide samples with various constituents in a number of matrices.

Reagent blanks should be analyzed to determine whether to evaluate contamination of supplies (reagents,
bottles, etc.) or carry over from one sample to the next. For example, in chromatographic analyses,
blanks may register non-zero values if the previous sample is not completely flushed from the column. A
minimum of 5% of samples analyzed should be reagent blanks.

Prior to analyzing unknown samples, your method should be calibrated using known standards. At least
three different dilutions of a known standard should be used and all reported values of unknowns should




                                                                                                          9
fall within the range of the calibration standards. If an unknown has a concentration higher than the
highest standard, the unknown should be diluted and re-tested.

Duplicate samples are analyzed in order to assess the precision of a method, which is calculated as a
percentage of the mean [ 100 x1  x2  / x ]. The acceptable limit for duplication is a function of the
chemical of interest and matrix, but is usually between 10%-20%. At least 5% of samples should be
measured in duplicate.

Detection levels are used to express the smallest detectable quantity for a certain procedure. Several
terms are used to estimate detection levels, which can lead to confusion. Four terms are most often used:
the instrument detection level (IDL), the lower level of detection (LLD), the method detection level
(MDL), and the level of quantitation (LOQ). The relationship among these levels is approximately
IDL:LLD:MDL:LOQ = 1:2:4:10. The practical detection limit (PDL) is used as a conservative limit for
cross-laboratory comparison and is approximately 3 – 5 times MDL.

Materials and Methods

1. Read the section in the laboratory manual on Laboratory Notebook and Grading.

2. Read the section in the beginning of the laboratory manual on Safety and Standard Practice. Copy
   down the emergency telephone number onto the first page of your laboratory notebook.

3. Prepare a floor plan of the laboratory indicating the locations of all furniture, exits, fire extinguishers,
   fire blankets, eyewash fountains, waste disposal containers (indicate types of waste), and telephones.
   Indicate the escape routes. This safety information should be present within the first few pages of
   your laboratory notebook.

Post-Lab Analysis

1. Go to the OSU Environmental Health and Safety web page (http://www.ehs.ohio-state.edu/) and
   download MSDS sheets for sulfuric acid and sodium hydroxide. Outline dangers and appropriate
   safety controls for use of each chemical.

2. Imagine a 9 lb. jug of concentrated sulfuric acid is dropped, breaks and spills acid on you and the
   floor. Describe in detail the exact sequence of steps you would take to rectify the situation.

3. Calculate the percent recovery of an analysis for arsenic if the measured concentration of the sample
   was 0.034 mg/L, the known addition was 0.10 mg/L, and the value of the sample with the known
   addition was 0.123 mg/L. The acceptable level of precision of known additions for arsenic is between
   80% and 120%. Do these data fall within the acceptable level of precision?

4. As a new environmental engineer/scientist, your company assigns you to collect and have analyzed
   samples from 15 groundwater wells at the Marion Engineering Depot. How many blanks and
   duplicates should you prepare and send to the lab along with your samples.

5. Calculate the precision of an analysis for the pesticide atrazine if two duplicates had measured
   concentrations of 23.0 ppb and 28.6 ppb. The acceptable level of precision of duplicates for
   pesticides is 20%. Do these data fall within the acceptable level of precision?

6. Calculate the MDL and PDL for mercury by atomic absorption spectroscopy if the IDL is 1 ng/L.




                                                                                                            10
                                           LABORATORY #2

                             SOLIDS CHARACTERISTICS OF WATER

Purpose

The purpose of this laboratory is to examine the turbidity, conductivity, and solids content of several
different water samples.

Background

Particulate matter is ubiquitous in all surface water and ground water systems, and can lend to water a
cloudy or hazy appearance. Due to their high specific surface area, suspended particles are efficient
adsorbents and they play a major role in regulating the transport and distribution of many chemical
compounds. Minute suspended particles may also indicate the presence of bacteria and non-bacterial
pathogens (e.g., Crytosporidium, Giardia, and viruses) and their presence inhibits disinfection processes.

The size spectrum of waterborne particles in both natural and polluted waters is continuous, spanning
roughly 0.001 m to 100 m (Figure 1). The composition of these particles can be organic (e.g., bacteria,
algae and viruses), inorganic (e.g., clay, sand, and iron oxides), or both. Suspended particles less than 1
m in diameter do not readily settle, and are considered colloids. This colloidal fraction may be highly
stable, hence mobile, and can significantly enhance the transport of adsorbed contaminants in surface and
ground water.

Turbidity. The concentration of solid particles in water can be measured based upon its turbidity using a
turbidimeter. This instrument measures the amount of scattered light in a water sample, and in general
scattering intensity increases with particle concentration. In this technique, a water sample in a glass
cuvette is irradiated with white light and a detector, located at an angle of 90 degrees from the incident
beam, measures the amount of scattered light. The preferred expression used to express turbidity is the
nephelometric turbidity unit (NTU). This nomenclature evolved from the original unit of turbidity
defined as the amount of scattered light originating from a 1 mg/L solution of SiO 2 of specific size. The
turbidity of a water in-situ can be estimated using a Secchi disk, and particle counters can be used to
provide quantitative information about the number of particles in a water or wastewater sample.

Direct microscopic examination provides information regarding the shape of particles and the degree of
particle aggregation. Although most theories of particle removal (for example, equations describing the
sedimentation of particles in settling basins) assume that particles are roughly spherical, particles can
possess a variety of shapes. Furthermore, particles are not often isolated units but instead exist as
members of large aggregates with densities that depend on the size of the aggregate. Particle aggregates
in a given aquatic system may be composed of organic and inorganic particles, polymeric material, as
well as metals and other constituents.

Gravimetric Analysis. A variety of separation techniques are commonly used to gravimetrically classify
solids by size and chemical characteristics (see Figure 2). Filtration is used to separate “suspended” or
“particulate” fractions from “dissolved” or “soluble” components. In this technique, glass-fiber filters are
used to remove particles as they pass through the deep mat of fibers by interception and impaction.
Because the average pore size and collection efficiency of commercial filters varies, it is always important
to specify the type of filter used, and pore size.

Evaporation separates water from dissolved and/or suspended matter. The drying temperature has an
important effect upon the results and weight losses due to volatilization of organic matter, water occluded


                                                                                                         11
                             VIRUS                             ALGAE


                                                     BACTERIA



          MOLECULES                                             SUSPENDED
                                                         (nonfilterable & settleable)

                                         COLLOIDS
                                         e.g., CLAYS

                                                  (1 micron)


      10-10        10-9        10-8        10-7        10-6        10-5       10-4        10-3
                                            Diameter (m)



 Figure 1. Size distribution of natural particles (after Stumm and Morgan, 1981).

in the interstices of crystals, water of crystallization, and gases from thermally induced decomposition
may occur. In addition, weight gain due to oxidation is possible. Consequently, two drying temperatures
are conventionally used: 103-105°C and/or 179-181°C. The lower temperature is used with samples con-
taining high concentrations of organic matter, which may undergo significant weight loss due to
volatilization and decomposition at the higher temperature. There is only slight decomposition of most
organic salts at 103°C. Some loss of CO2 can be expected from the conversion of bicarbonate to
carbonate during the dehydration process. Occluded or bound water is not completely removed at 103°C;
however, its removal is virtually complete at 180°C. At 180°C, thermal decomposition of ammonium
salts (especially ammonium carbonate) may occur. The type of cation in the salt greatly affects the
degree of decomposition at a given temperature.

Solids are operationally characterized as either volatile or nonvolatile (in some texts the term “fixed” is
used in place of nonvolatile). Volatile solids are those that volatilize when samples are heated to a
temperature of 550°C. The volatile fraction is primarily composed of organics, whereas the fixed fraction
is mostly inorganic. To measure these quantitites, samples are placed in a “muffle” oven at 550°C for a
fixed period of time. The remaining sample is weighed again and the mass loss represents the volatile
fraction. Special filters, made of noncombustible glass fibers, must be used.

Conductivity. The conductivity of a solution is a measure of its ability to carry an electric current. It
varies with temperature and depends on the presence of ions and their total concentration, mobility, and
valence. In practice, conductivity can be used as a measure of the dissolved solids in water, and the
greater the dissolved solids the greater the ability of the water to carry electric current. Conductivity may
also be referred to as specific conductance. The conductivity of deionized water is typically between 0.5
and 3.0 S/cm, and that of potable water ranges between 50 and 1500 S/cm. Wastewater conductivity
may be as high as 10,000 S/cm.


                                                                                                          12
                                        Sample



                                  0.45 m filtration


                 Heating                                    Heating                                Heating
                 at 105 C                                  at 105 C                              at 105 C



           Total Dissolved                             Total Suspended                            Total Solids
            Solids (TDS)                                Solids (TSS)                                 (TS)


                Heating at                                  Heating at                            Heating at
                 550 C                                      550 C                                550 C




     Volatile                Nonvolatile         Volatile                Nonvolatile    Total                  Nonvolatile
    Dissolved                Dissolved          Suspended                Suspended     Volatile                  Total
     Solids                    Solids             Solids                   Solids      Solids                    Solids
     (VDS)                                        (VSS)                                (TVS)



   Figure 2. Operational definition of different types of solids.
A number of units for conductivity are in current use. The tradition unit for conductivity is 1/ohm-cm,
or mho/cm. In the SI system of units, siemens (S) is the reciprocal of ohm and conductivity is often
reported in units of S/cm. Thus, S/cm and mho/cm are equivalent units. The total dissolved solids
(in mg/L) of a water sample can be estimated by multiplying the conductivity (in S/cm) by an empirical
constant (usually between 0.55 and 0.90). This empirical constant should be determined for a particular
water sample by comparing conductivity measurements to a direct gravimetric analysis.

Materials and Methods

In this laboratory, you will work in groups of 4-6 people. The following glassware and reagents are
required for the analysis of turbidity, conductivity, and solids content.

      Glassware. (250mL volumetric flask [1 per analysis team], 300mL beaker, 100 mL
       beaker, 25mL volumetric pipette.
      Check standard. (prepared by the TA by suspending 100 mg/L kaolin in a solution of
       100 mg/L NaCl in deionized water)
      Turbidity meter and sample cells.
      Filtration apparatus. (vacuum flask, clamp, 2-piece membrane holder, and pump)
      Filter papers. (rinsed, dried at 103 – 105 C, and stored in desiccator by TA prior to lab)
      Aluminum weigh boats. (ignited at 550 C by TA prior to lab)



                                                                                                                      13
       Tweezers.
       Conductivity meter and probe.

Determine the turbidity, conductivity, and solids content (e.g., TSS, TDS & TS) of tap water, Olentangy
river water, Scioto river water, one blank sample (deionized water), one check standard (100 mg/L
kaolin+100 mg/L NaCl), and a duplicate as follows:

Turbidity

To determine the turbidity of a water sample, carry out the following:

1. Obtain 200 mL of an unknown water sample from the source provided and place it in a 250-300 mL
   beaker. Be sure that the source is well mixed prior to obtaining your sample.

2. Turn the turbidimeter on and allow it to warm up for approximately 5 minutes.

3. Using the standard turbidity samples, verify that the turbidimeter is working properly. Be sure to
   mark down the standard turbidity indicated on the sample, as well as the turbidity as measured by the
   instrument.

4. Place 25mL of unknown sample into a clean, glass turbidity tube and insert the sample in the
   turbidimeter.

5. Measure the turbidity to the nearest 0.1 NTU.

Gravimetric Solids Determination

In this portion of the lab, your group will determine the TS and TSS for each sample using a gravimetric
approach. For each sample, carry out the following:

Total Solids (TS)

1. Zero the balance.

2. Weigh an ignited and cooled aluminum evaporation dish to the nearest 0.1mg.

3. Add 25.0 mL (or 50 mL if large dishes available) of sample solution using a volumetric pipette.

4. Evaporate to dryness in an oven at 103°C. It is useful to place your samples in the oven on a piece of
   paper identifying your group and the contents of the sample.

5. Cool in a desiccator to room temperature (Choose 1-2 group members to place dishes in the
   desiccator after one day of drying).

6. Zero the balance (Choose 1-2 group members to weigh the dishes after one day in the desiccator).

7. Remove and reweigh the aluminum evaporation dishes to the nearest 0.1mg. Make sure to remove
   and weigh each dish one at a time. Use a consistent time interval between removal and weighing
   (e.g., 1 min).




                                                                                                       14
Total Suspended Solids (TSS)

1. Zero the balance.

2. With tweezers, transfer the pre-washed and dried filter from the desiccator to a balance and determine
   the weight to the nearest 0.1mg.

3. Place the filter in a filtration apparatus.

4. Pour through exactly 0.250 liter of sample measured using a volumetric flask and apply suction.

5. Wash the sample container with 50 mL of Deionized water, then the filter holder and filter with two
   10mL portions of Deionized water or until all of the solids are contained on the filter. All wash water
   should be passed through the filter. After all the sample and rinse water has passed through the filter
   discard the solution in the filter flask.

6. Remove the filter with tweezers and evaporate to dryness at 103°C (use an un-dried aluminum dish to
   hold filter in drying oven and to provide identification of filter).

7. Cool in the desiccator (Choose 1-2 group members to place filters in the desiccator after one day of
   drying).

8. Zero the balance (Choose 1-2 group members to weigh the dishes after one day in the desiccator).

9. Remove and reweigh each filter to the nearest 0.1mg. Make sure to remove and weigh each filter one
   at a time. Use a consistent time interval between removal and weighing (e.g., 1 min).

Conductivity

1. Turn on the conductivity meter.

2. Place the conductivity probe in the conductivity standard solution and verify the meter is responding
   accurately.

3. Obtain approximately 200 mL of sample and place in a 250 mL beaker.

4. Place the probe in the sample solution and measure the conductivity. Record your data in S/cm.

5. Estimate the TDS (in mg/L) by multiplying the conductivity by a conversion factor of 0.55.




                                                                                                       15
Pre-Lab Questions

1. A water sample has a specific conductance of 235 S/cm. Assume a conversion factor and calculate
   the TDS in mg/L.

2. Calculate the concentrations of TS, TSS, TDS, TVS, VSS and VDS in mg/L based on the gravimetric
   analysis provided below. Present your results in a tabular format.

In the first experiment, 250.0mL of sample was placed in an evaporation dish and heated at 103°C for
about 6 hours. Next, the sample was placed in a muffle oven and heated again, this time at 550°C. The
following data was collected:

                        sample size :                                250.0 mL
                        Tare of dish =                               2.3315 g
                        dish + sample after heating at 103°C =       2.4008 g
                        dish + sample after heating at 550°C =       2.3545 g

In the second experiment, 500.0mL of sample was filtered onto a glass fiber filter and then the filter was
dried at 105 and 550°C. The data for this experiment is given below.

                        sample size :                                500.0 mL
                        Tare of filter =                             1.3255 g
                        filter + sample after heating at 105°C =     1.4170 g
                        filter + sample after heating at 550°C =     1.3601 g

In the last experiment, the filtrate from the second experiment was analyzed. Here, 250.0mL of filtrate
was placed in an evaporation dish and then heated and weighed. The results are:

                        sample size :                                250.0 mL
                        Tare of dish =                               2.3324 g
                        dish + sample after heating at 105°C =       2.3559 g
                        dish + sample after heating at 550°C =       2.3382 g


Post-Lab Analysis

1. Calculate TS, TSS, and TDS based on your gravimetric analysis.

2. Prepare a summary table of your data similar to the one shown below:

Sample             Turbidity     Conductivity        TDS             TS          TSS         TDS
                    (NTU)          (S/cm)         (C x 0.55)      (mg/L)       (mg/L)     (TS-TSS)
                                                    (mg/L)                                  (mg/L)
Blank
Check Standard
Tap
Scioto
Olentangy-1
Olentangy -2




                                                                                                          16
3. Calculate the % recovery of your check standard for TDS (based on conductivity), TDS (based on
   gravimetric analysis), TS, and TSS. Are they acceptable? If not, what errors may have been
   introduced into your analysis to account for this?

4. What is the precision of all your analyses based on your duplicate samples?

5. How do your measurements of TDS based on conductivity and the gravimetric analysis compare?
   What sources of error may contribute to any differences.

6. How does your data for tap water compare to values obtained by the City of Columbus for the Dublin
   Road and Hap Creamen water treatment plants (see p. 8)? What factors may contribute to differences
   in the data?

7. How does your data for Scioto river water compare to values obtained by the City of Columbus for
   the influent to the Dublin Road water treatment plant? What factors may contribute to differences in
   the data?




                                                                                                      17
                                           LABORATORY #3

                               ACID AND BASE STANDARDIZATION

Purpose

The objective of this laboratory is to prepare and standardize stock solutions of acid (sulfuric) and base
(sodium hydroxide).

Background and Theory

Acid Standardization. Standardization of acid solutions is accomplished by precisely measuring the
volume of acid needed to neutralize a known amount of base. The usual base primary standard is sodium
carbonate, however, tris-(hydroxymethyl)aminomethane, known also as “TRIS” or “THAM” is also used.
The chemical formula for TRIS is (HOCH2)3CNH2, and it has a molecular weight of 121.135. TRIS
reacts with acids by adding one proton to the amino group, becoming (HOCH 2)3CNH3+. Thus, its
equivalent weight is also 121.135. The proton can come from either the surrounding water or from the
added acid.

TRIS is available commercially in primary-standard purity, and its rather large formula weight (121.135)
minimizes weighing errors. TRIS is hygroscopic and prior to use must be dried at 105°C so that
absorption of atmospheric moisture does not impair weighing. The equivalence point of TRIS lies
between pH 4 and 5 and therefore bromcresol green (which changes color from blue to yellow at pH 4.5)
is a good choice of indicator.

Base Standardization. Bases are standardized by titrating precisely weighed quantities of acid. The usual
choice for primary acid standard is potassium acid phthalate, also known as potassium hydrogen phthalate
or “KHP”. The formula for KHP is KHC8H4CO4, and its molecular weight is 204.223. It only gives up
one proton during titration, so its equivalent weight is also 204.223.

KHP is dried at 105°C (but not higher) and cooled in a desiccator. The resulting solid is nonhygroscopic
and has a relatively high molecular weight. The equivalence point occurs when all the KHP has been
stripped of protons. The pH at the equivalence point occurs between 8 and 10 and, therefore,
phenolphthalein is a good indicator.

Significant Digits. In any analytical procedure it is important to consider significant digits, in particular
when specifying the concentration of standard solutions. In general, all reported digits should be
definitively known, except the last digit, which may be in doubt. For example, an analyst reports the
concentration of TSS as 76.5 mg/L. The analyst is confident about the 76, but is not sure if it is 0.5, 0.7
or 0.4, etc. When values for a parameter (e.g., TSS) are presented in a column for multiple samples, all
values in the column need not have the same number of significant digits.

The standard deviation of an analysis is used to determine which digits are significant. For example, the
TDS of a sample measured as 1467 mg/L, with a standard deviation of 40 mg/L, should be reported as
1470 mg/L. If the standard deviation is 100 mg/L, report the measurement as 1500 mg/L.

When multiplying or dividing, the digits in the final answer should equal those in the factor with the
fewest significant figures. If adding or subtracting, the number with the fewest decimal places, not the
number with the fewest significant figures, puts the limit on the number of places that may be justifiably
carried in the sum or difference.




                                                                                                          18
Materials and Methods

In this laboratory you will work in groups of 4-6 people. The following glassware and reagents are used
in the laboratory:

       Glassware. (1 L volumetric flask, 10 or 25mL Mohr pipette, 1 L glass stock acid
        bottle, 500mL Erlenmeyer flask (3), 50mL burette, 1 L polyethylene stock NaOH
        bottle)
       Concentrated sulfuric acid (H2SO4). 17.6M (35.2N), molecular weight of 98.079.
       Sodium hydroxide (NaOH). molecular weight of 39.997.
       Tris-(hydroxymethyl)aminomethane (TRIS). molecular weight of 121.135.
       Potassium hydrogen phthalate (KHP). molecular weight of 204.223.
       Bromcresol Green. prepared by the TA by dissolving 0.1 g bromcresol green into
        100mL deionized water and adding 3 drops of 0.2 N NaOH.
       Phenolphthalein. prepared by the TA by dissolving 0.2 g phenolphthalein into
        200mL deionized water and 200mL 95% ethanol.

Acid Standardization. Prepare stock 0.2 N solution of sulfuric acid and standardize with TRIS. To
prepare your solution, follow the procedure below.

1. Calculate the volume of concentrated H2SO4 needed to produce 1 L of 0.2N H2SO4.

2. Put about 0.5 L of deionized water into a 1 L volumetric flask.

3. Put a suction bulb (Propipette) on a graduated (Mohr) pipette, being careful to lubricate the end of the
   pipette with deionized water and using a gentle, axial, rotating force to seat the bulb.

4. While working in the fume hood, draw the needed amount of concentrated H2SO4 into the pipette
   and drain the pipette into the volumetric flask. Be sure you are proficient using the pipette prior to
   attempting the procedure with acid. (If not proficient, then practice with deionized water.)

5. Carefully fill the volumetric flask to the mark with deionized water and mix by holding the flask in
   two hands (one under the base and one on the neck). Mix thoroughly by gently swirling the contents
   and turning the flask end-over-end.

6. Drain the volumetric flask into a 1 L glass stock solution bottle, and label the bottle with (a) your
   name, (b) the date, (c) the solution it contains, and (d) the approximate concentration of the solution.
   All solutions should be labeled in this way.

To standardize the 0.2 N sulfuric acid solution prepared above, do the following:

1. Dry 3 to 5 g TRIS at 105°C and cool in a desiccator. (This has been done for you by the TA to save
   time.)

2. Weigh out three samples of TRIS between 0.7 and 0.75 g, and add to three separate 500mL
   Erlenmeyer flask. Record the exact weight of each TRIS sample to 4 significant figures.

3. Add approximately 150mL of deionized water to each flask to dissolve the TRIS.

4. Add 2 to 3 drops of bromcresol green indicator to each flask.


                                                                                                        19
5. Fill a clean 50mL burette with the stock 0.2 N H2SO4 solution.

6. Titrate the TRIS in each Erlenmeyer flask until the solution turns yellow by adding small volumes of
   acid and swirling the flask to mix. The endpoint should be about 30mL, and as you approach the end-
   point, add the acid drop-wise. The equivalence point determined using your indicator should be
   accurate to the nearest drop. Record the volume of acid dispensed from the burette to the nearest
   0.01mL.

7. Calculate the normality of the stock acid solution assuming the TRIS is 100 percent pure.

Base Standardization. Prepare a stock 0.2 N solution of sodium hydroxide and standardize your stock
solution with KHP. To prepare your base solution, follow the procedure below.

1. Calculate the weight of sodium hydroxide pellets needed to produce 1 L of 0.2 N NaOH.

2. Select and zero a plastic-weighing dish or weighing paper.

3. Weight out the amount of NaOH needed to prepare 1 L of 0.2 N NaOH solution by adding small
   amounts of the pellets to the plastic-weighing dish or paper. Use a scoopula. You can make this
   weighing process go more quickly by first determining the weight of an average pellet and then
   counting out the approximate number of pellets needed. Remove and dispose of pellets that may
   have fallen on the scale. NaOH is corrosive and will quickly ruin the instrument.

4. Put about 0.5 L of deionized water into a 1 L volumetric flask.

5. Carefully add the pellets to the volumetric flask and let the pellets dissolve. You can hasten this by
   holding the flask in two hands and gently swirling the contents.

6. When the pellets have dissolved, carefully fill the volumetric flask to the mark with deionized water.
   Mix by inverting the flask several times after securing a cap.

7. Drain the volumetric flask into a 1 L polyethylene stock solution bottle and label the bottle with (a)
   your name, (b) the date, (c) the solution it contains, and (d) the approximate concentration of the
   solution.

To standardize the 0.2 N sodium hydroxide solution prepared above, do the following:

1. Dry 3 and 5g KHP at 105°C and cool in a desiccator. (To save time, this has been done for you by
   the TA.)

2. Weigh out three (3) quantities of KHP, each between 1.1 and 1.3g, and add each quantity to a
   separate 500mL Erlenmeyer flask. Be sure to record the exact weight of each quantity of KHP to 4
   significant figures.

3. Add approximately 150mL of deionized water to each flask to dissolve the KHP.

4. Add 2 to 3 drops of phenolphthalein indicator to each flask.

5. Fill a clean 50mL burette with the stock 0.2 M NaOH solution.

6. Titrate the KHP in each Erlenmeyer flask until the solution turns pink. Titrate by adding small
   volumes of base and swirling to mix. The endpoint should be about 30mL. As you approach the end-


                                                                                                      20
    point, add the base drop-wise. The equivalence point determined using your indicator should be
    accurate to the nearest drop. Record the volume of base added to the nearest 0.01mL.

7. Calculate the normality of the stock base solution assuming KHP is 100 percent pure.

Pre-Lab Questions

1. Calculate the volume of concentrated H2SO4 needed to produce 1 L of 0.2N H2SO4 and the weight of
   sodium hydroxide pellets needed to produce 1 L of 0.2 N NaOH.

2. Suppose the following calculation must be made to obtain the result of your standardization. What is
   the answer to the computation (in eq/L) with the correct number of significant digits?

         1.1042 g 1000 mL / L
   N
        30 .40 mL  204 .223 g / eq

3. The iron concentration in AMD (Acid Mine Drainage) is determined to be 1.034 mg/L. If the
   standard deviation of the analysis is 0.010 mg/L what is the iron concentration with the correct
   number of significant digits?

4. The acid standardization procedure yields a titrant volume of about 30mL. Why is this desirable? A
   volume of 5 or 10mL would be much faster and would consume less material. Why would not 5 or
   10mL suffice?

Post-Lab Analysis

1. Calculate the average normality and standard deviation of your acid and base standardizations and
   record the results in your lab notebook with the correct number of significant digits. Use these values
   for all subsequent analyses using these standardized solutions. Re-label your stock acid and stock
   base bottles with the average normality and standard deviation.




                                                                                                       21
                                           LABORATORY #4

                           DETERMINATION OF pH AND ALKALINITY

Purpose

The objective of this laboratory is to determine which chemical species comprise the alkalinity of a
sample and to produce a pH-titrant volume curve for the titration of natural alkalinity with a strong acid.

Background and Theory

pH is used to express the activity of the hydrogen ion in solution and is defined as the negative logarithm
of the hydrogen-ion activity (i.e., pH = -logH+). Relatively high hydrogen-ion activities have low pH
values and are “acidic.” Conversely, at low hydrogen-ion activity the pH is high and the solution is
“alkaline.” pH is commonly referred to as a “master variable” because it controls a number of chemical
processes, including speciation, complexation, precipitation and dissolution. The most convenient
method of determining pH involves measuring the electrical potential that forms across a membrane that
separates two solutions of differing pH. The greater the difference in pH of the two solutions, the higher
is the potential. Thus, if the pH of the reference solution is known, the sample pH can be determined
based on the potential formed. Numerous electrodes commercially available for pH measurement are
based upon this principle, and incorporate a glass-membrane in the electrode tip.

Alkalinity is a measure of the capacity of a water sample to neutralize strong acid. In natural waters this
capacity is attributable to bases such as HCO3-, CO32-, and OH- as well as to species often present in small
concentrations such as silicates, borates, ammonia, phosphates, and organic bases. Alkalinity in natural
waters affects a wide range of processes such as coagulation in water treatment operations, buffering
capacity of lakes and rivers, and ammonia stripping, to name a few.

To determine the total alkalinity, a known volume of sample is titrated with a standard solution of a
strong acid to a final pH value of approximately 4.5. This endpoint is commonly indicated by the color
change of the indicator1 bromcresol green, and it corresponds to the pH at which the conversion of
bicarbonate to carbonic acid is complete. The H+ added is the stoichiometric amount required for the
following reactions:

        H+ + HCO 3-  H2 CO3                                                                           [1]

At the pH corresponding to the endpoint of the titration for total alkalinity the solution contains only
H2CO3 and H2O.

For solutions with a high starting pH (nominally greater than 8.3), the alkalinity is determined in two
steps. In the first step, the titration is conducted to an endpoint pH of approximately 8.3 (see Figure 3),
where the conversion of carbonate to bicarbonate is complete. At this endpoint, a stoichiometric amount
of H+ has been added to complete the following two reactions:

        H+ + OH -  H2O                                                                                [2]
        H+ + CO3 2-  HCO3-                                                                            [3]
1
 Indicators for acid-base titrations are highly colored weak acids or bases that undergo substantial color
shifts when changed from the neutral to the ionized form. A number of indicators are useful over a range
of pH values and when choosing the appropriate indicator for an acid-base titration it is important to
match the pKa of the indicator to the pH of the endpoint of the titration.


                                                                                                         22
       12
       11
       10                                               Point of
                                                        Inflection
       9
       8
  pH




       7
       6                                                                   Point of
                                                                           Inflection
       5
                    +      -
                   H + OH  H2O                     +           -
       4                                       H + HCO3 
                    +      2-
                   H + CO3  HCO3
                                       -       H2CO3
       3
       2
            3.00    2.50       2.00   1.50   1.00        0.50       0.00   -0.50        -1.00   -1.50     -2.00

                                                mL Acid


Figure 3. Titration curve for the carbonate system.

This is often termed the phenolphthalein endpoint, based on the use of the indicator phenolphthalein (pKa
~ 8.3). This endpoint is typically poorly defined. In the second step of the titration, the solution is titrated
to convert bicarbonate to carbonic acid as previously described and shown in equation [1].

The expressions in Table 1 can be used to predict contributions from the three principal forms of
alkalinity (CO32-, HCO3- and OH-), based on the phenolphthalein (P) and total alkalinity (T) endpoints.
The values of P and T in Table 1 should be expressed in terms of meq/L or mg/L calcium carbonate.

Table 1. Forms of Alkalinity

                           Titration         Hydroxide                     Carbonate                    Bicarbonate
                            Result           Alkalinity                    Alkalinity                    Alkalinity
                             P=0                 0                             0                             T

                              1                         0                          2P                      T-2P
                           P T
                              2
                              1                         0                          2P                         0
                           P T
                              2
                              1                     2P-T                     2(T-P)                           0
                           P T
                              2
                            P=T                         T                          0                          0

Note that the true endpoint for alkalinity determinations is slightly variable and is best determined by
measuring the pH at different points of the titration and making a plot of this versus volume of added
acid. The true endpoints are then determined from this plot by the points of inflection.


                                                                                                                      23
Materials and Methods

In this laboratory, you will work in groups of 4-6 people. The following glassware and reagents
(prepared prior to laboratory) are used in the analysis of alkalinity.

       Glassware. (1 L volumetric flask, 25mL Mohr pipette, 100mL volumetric flask, 500mL
        Erlenmeyer flask or beaker, 50mL burette)
       0.2 N sulfuric acid.
       Phenolphthalein indicator.
       Bromcresol green indicator.
       Alkalinity Standard. (100 mg/L as CaCO3; Prepared by dissolving 103 mg Na2CO3 in 1
        L of Deionized water)
       pH meter and electrode.
       pH 4, 7 and 10 buffer solutions.

Prepare the following:

          0.02 N sulfuric acid. Calculate the volume of 0.2 N H2SO4 needed to produce
          1 L of 0.02 N H2SO4. Put about 0.5 L of deionized water into a 1 L volumetric
          flask. Using a graduated Mohr pipette, draw the needed amount of 0.2 N H 2SO4
          into the pipette and drain the pipette into the volumetric flask. Carefully fill the
          flask to the mark with deionized water and gently swirl the contents. Label the
          volumetric flask with your name, the date, the solution it contains, and the
          approximate concentration.

Determine the alkalinity for tap water, Scioto river water, Olentangy river water, one blank, one check
standard, and one duplicate using appropriate indicators. For one of the river water samples and the
check standard also use direct pH measurement to determine alkalinity. Perform the following steps for
each of the unknown samples.

1. Examine and clean the pH electrodes and check and make sure it is filled with electrolyte. Note: pH
   probes are fragile and should be handled with care.

2. Calibrate the pH meter using the standard pH4, pH7 and pH10 buffer solutions. Note: the pH meter
   should be in “stand-by” mode when the electrode is not submerged in either buffer or acid
   solution.

3. Fill a 50mL burette with your 0.02 N acid solution.

4. Using a volumetric flask, add 100mL of sample to a 500mL Erlenmeyer flask (use a beaker for
   determining alkalinity by pH measurement).

5. Clean the electrode with deionized water, put the electrode into the water sample and measure the
   initial sample pH.

6. If the pH of the sample is greater than the phenolphthalein endpoint add 2 to 4 drops of the indicator
   and the solution will turn pink/red. If not, then go to step 8.

7. Titrate the alkalinity of the sample to the phenolphthalein endpoint where the solution is colorless.
   Carefully record the acid volume (and pH) after approximately every 0.2 – 0.4 mL addition (use
   smaller increments near inflection point). Make note of the pH that you observe the color change.


                                                                                                      24
8. Add 2 to 4 drops of bromcresol green indicator. The solution should be blue.

9. Titrate the sample to the bromcresol green endpoint (a faint yellow) and record the total volume of
   acid added (and pH) at approximately 0.2 – 0.4 mL intervals. This endpoint gives the total alkalinity.
   Make note of the pH that you observe the color change.

Pre-Lab Questions

1. What is the dominant carbonate species at the phenolphthalein endpoint? Demonstrate your answer
   by numerical calculation. (Hint: look up the equilibrium for the carbonate system from the text.)

2. Why is bromcresol green a good indicator for total alkalinity?

3. Show for a 0.02 N H2SO4 solution that 1 mL of titrant is equivalent to 1 mg of alkalinity as CaCO3.

4. Calculate the total alkalinity (in mg/L as CaCO3) of a sample if 23.4 mL of 0.02 N H2SO4 are
   required to reach the bromcresol green endpoint and a total of 150 mL of sample was used.

5. Calculate the hydroxide, carbonate and bicarbonate contributions to alkalinity (in eq/L) if 5.5 mL and
   25.5 mL of 0.02 N H2SO4 are required to reach the phenolphthalein and bromcresol endpoints,
   respectively. Sample size is 100 mL.

Post-Lab Analysis

1. On a single graph, plot pH versus total volume of acid added for the one river water sample and check
   standard you recorded pH-volume data. Mark the values at which you observe a color change. How
   does the defined endpoint pH values (4.5 and 8.3) compare to the pH values in which you observed a
   color change? How might any differences influence your alkalinity calculations?

2. Calculate the phenolphthalein and total alkalinity in units of mg/L calcium carbonate using indicator
   data as well as actual pH measurements.

3. Prepare a table similar to the one below summarizing your data:


               Sample                   pH             Carbonate             Total
                                                       Alkalinity          Alkalinity
                                                     (mg/L CaCO3)        (mg/L CaCO3)
               Blank
               Check Standard
               Tap
               Scioto
               Olentangy
               (Replicate)



4. Calculate the % recovery of your check standard. Is it acceptable? If not, what errors may have been
   introduced into your analyses to account for this?

5. What is the precision of your analyses based on your duplicate samples?



                                                                                                         25
6. How does your data for tap water compare to values obtained by the City of Columbus for the Dublin
   Road and Hap Creamen water treatment plants? What factors may contribute to differences in the
   data?

7. How does your data for Scioto river water compare to values obtained by the City of Columbus for
   the influent to the Dublin Road water treatment plant? What factors may contribute to differences in
   the data?

8. Use Table 1 to estimate the original concentrations of hydroxide, carbonate, and bicarbonate in the
   sample. Convert the species concentrations to moles per liter. What assumptions are implied in these
   calculations?




                                                                                                    26
                                            LABORATORY #5

                             TITRIMETRIC METHOD FOR HARDNESS

Purpose

The purpose of this laboratory is to determine hardness in a number of water samples using the method of
chelation.

Background and Theory

Water hardness is a means of quantifying the concentration of multivalent cations. At sufficiently high
concentration, hardness produces undesirable effects that include the loss of detergency with the use of
soaps, the formation of precipitates or “scale” in boilers, hot-water heaters, and distribution pipelines, and
                                2+          2+                                           2+      2+         2+
undesirable taste. While Mg and Ca are the primary hardness cations, Fe , Mn , and Sr
contribute to water hardness in certain cases.

The total hardness of water is defined as the summation of the concentrations for hardness cations, or
“hardness” [Ca2+] + [Mg2+]. Hardness units are generally expressed in mol/vol., equiv./vol., or as the
equivalent mass of CaCO3/vol. Equivalent concentrations in this case are computed on the basis of the
metal ion charge. The total hardness may be divided into two forms; “carbonate” and “non-carbonate”
hardness. When the hardness is numerically greater than the alkalinity due to carbonate and bicarbonate,
that amount of hardness equivalent to the total alkalinity (assuming OH-alkalinity is negligible) is called
carbonate hardness. The amount of hardness in excess is called the non-carbonate hardness. When the
hardness is numerically equal to or less than the sum of carbonate and bicarbonate alkalinity, all hardness
is carbonate hardness.

Most frequently, hardness is determined analytically by complex-formation titrations, which involves
adding a complexing or chelating agent to form coordination compounds, or complex ions, with metal
cations. Chelates are produced when a metal ion coordinates with two or more donor groups of a single
complexing molecule. In general, the stability of a chelate is related in part to the number of coordinative
bonds that can form between the chelating agent and the metal ion.

One of the most common hardness test methods involves the aminopolycarboxylic chelating agent,
ethylene diamine (N,N,N’,N’) tetraacetic acid (called EDTA). The EDTA molecule has the following
structure:

            H2                              H2
HOOC        C                               C  COOH
                  N    H2     C       N
                       C      H2
           C                                C
HOOC       H2                               H 2 COOH

The EDTA molecule has six potential sites for bonding with a metal ion: the four carboxyl groups
(pKa1=2.0, pKa2=2.8, pKa3=6.2, pKa4=10.3) and the two amino groups. Metal ions usually have a
coordination number of six, and thus form very stable complexes with EDTA. If EDTA is not fully
deprotonated, the complexes are less stable.

In the determination of hardness, samples are titrated with EDTA. Several competing equilibria are
involved. The sample solution is buffered at pH 10±0.1 as a compromise between chelate stability and
the prevention of metal ion precipitation (especially CaCO3 or Mg(OH)2). An ammonia buffer is used


                                                                                                           27
since ammonia forms weak complexes with hardness cations and helps to prevent their precipitation.
EDTA and EDTA-metal complexes are not colored, and an additional chelating agent and indicator,
calmagite, is used to determine the endpoint of the EDTA titration. Calmagite is deep red when
complexed with divalent ions, and can be detected visually. The ammonia buffer is usually spiked with a
trace amount of Mg2+-EDTA to facilitate endpoint detection in the event the sample does not contain
Mg2+.

As the EDTA titrant is added to the sample solution, it reacts first with Ca 2+, then with Mg2+, since the
Ca2+-EDTA complex is more stable than the Mg2+-EDTA complex:

    EDTA+ Ca2+  Ca2+-EDTA                       K=1010.7                                            [2a]
    EDTA+ Mg2+  Mg2+-EDTA                       K=108.7                                             [2b]

Metal ions are also readily extracted from ammonia complexes because metal-EDTA complexes are
much more stable. After adding sufficient EDTA, such that all of the free Mg 2 is complexed, further
addition of EDTA removes Mg2+ from the Mg-calmagite complex, causing the indicator to change to a
blue color.

    Mg-calmagite(red) + EDTA  Mg-EDTA + calmagite(blue)                                               [3]

The resulting color change allows for the determination of the endpoint of the titration. Since one mole of
EDTA binds with exactly one mole of Ca 2+ or one mole of Mg2+, the total moles of EDTA added
represents the total moles of hardness, or

                                moles EDTA titrated 1 mole CaCO3 100, 000mgCaCO 3
Hardnes s(mg / L CaCO 3 )                          x            x                                     [4]
                                 vol.sample titrated 1 mole EDTA    mole CaCO 3

Materials and Methods

In this laboratory, you will work in groups of 4-6 people.

      Glassware. (25mL volumetric pipette, 100mL volumetric flask, 50mL burette, plastic
        beakers).

Prepare the following stock solutions (the buffer, indicator, EDTA, and 1000 mg/L stock standard
calcium solution have been prepared for you. Each group is responsible for preparing the 250 mg/L
calcium check standard):

       Buffer Solution: Store in a plastic container. Dissolve 16.9g ammonium chloride,
        NH4Cl, in 143mL concentrated ammonium hydroxide, NH4OH; add 1.25g of magnesium
        salt of EDTA and dilute to 250mL with deionized water. This solution should be kept in
        the fume hood.

       Indicator Solution (Calmagite 0.1%): The dye calmagite is 2-Hydroxy-1-(2-hydroxy-
        5-methylphenylazo)-4-naphthalenesulfonic acid. To prepare a 0.1% solution, dissolve
        0.1g of calmagite in 100mL of water. This indicator will deteriorate so it will need to be
        made fresh each year.

       Standard EDTA titrant, 0.01 M: Store in Polyethylene or Pyrex bottles. Dissolve
        3.723g analytical reagent-grade disodium ethylenediamine tetraacetate dihydrate, also



                                                                                                        28
        called (ethylenedinitrilo)-tetraacetic acid disodium salt (EDTA), Na2H2C10 H12O8N2
        *2H2O, in one liter deionized water. Standardize against the following calcium solution.

       Stock Standard Calcium Solution (1000 mg CaCO3/L): Place 1.000g anhydrous
        calcium carbonate, CaCO3, powder (primary standard or special reagent low in heavy
        metals, alkalis, and magnesium) in a 500mL Erlenmeyer flask. Place a funnel in the neck
        of the flask and add, a little at a time, 1+1 HCl (1+1 indicates one volume of reagent HCl
        is diluted with 1 equal volume of deionized water) until all the CaCO3 has dissolved.
        Add 200mL deionized water and boil for a few minutes to expel CO2. Cool, add a few
        drops of methyl red indicator (0.1% solution), and adjust to the intermediate orange color
        by adding 3N NH4OH or 1+1 HCl, as required. Transfer this to a one liter volumetric
        flask and fill to the mark with deionized water.

       Check Standard (250 mg CaCO3/L): Using a 25 mL volumetric pipette, transfer
        exactly 25 mL of stock standard calcium solution to a 100 mL volumetric flask. Fill to
        the mark with Deionized water.

Determine the hardness of tap water, Scioto river water, Olentangy river water, one blank, one check
standard, and one duplicate using the EDTA titrimetric approach as follows:

1. Dilute 25.0mL sample to about 50mL with deionized water in a plastic beaker. Add 1 to 2mL buffer
   solution (usually 1mL will be sufficient to give pH of 10.0 to 10.1). Note: Do not take pure buffer
   solution outside of the fume hood onto your bench. Add the buffer to your sample and
   immediately perform your titration (to minimize hydroxide from evaporating and prevent the
   precipitation of calcium carbonate). Once the titration is complete, pour the sample down the
   sink and flush the sink with generous quantities of water.

2. Add 1 to 2 drops of indicator solution (calmagite). Then titrate slowly, with continuous stirring, until
   the last reddish tinge disappears, adding the last few drops at 3 to 5 second intervals.

3. The solution should be blue at the endpoint.

4. Calculate and report the hardness of your samples in mg/L as CaCO3. The hardness is calculated as:

                                               A x B x 1000
        Hardness (EDTA) as mg/l CaCO3 =
                                                    C

    where A is the titrant volume (mL), B is the conversion factor for calculating mg of CaCO3
    equivalent to 1.00mL of the EDTA titrant, and C is the sample volume (mL).

Pre-Lab Questions

1. Show that if the standard EDTA titrant is 0.01M, then 1mL of titrant is equivalent to 1mg hardness as
   CaCO3 (This is parameter “B” in the equation above).

2. Calculate the total hardness of a sample if 10 mL of 0.01 M EDTA is required to reach the calmagite
   endpoint and 25 mL of sample is used in the analysis.

3. Calculate the carbonate and non-carbonate hardness of a sample if the total alkalinity is 130 mg/L as
   CaCO3 and the total hardness is 175 mg/L as CaCO3.



                                                                                                        29
Post Lab Analysis

1. Based on your alkalinity analysis last week and your total hardness data this week, calculate the
   carbonate and non-carbonate hardness for Scioto and Olentangy river waters in mg/L as CaCO3.

2. Do the same for the tap and blank samples and summarize your data as shown below:


          Sample                Total Hardness      Carbonate      Non-Carbonate
                                  Hardness          Hardness         Hardness
                                (mg/L CaCO3)          (mg/L        (mg/L CaCO3)
                                                     CaCO3)
          Blank
          Tap
          Scioto
          Olentangy
          (Duplicate)



3. Calculate the % recovery of your check standard. Is it acceptable? If not, what errors may have been
   introduced into your analysis to account for this?

4. What is the precision of your analyses based on your duplicate samples?

5. How does your data for tap water compare to values obtained by the City of Columbus for the Dublin
   Road and Hap Creamen water treatment plants? What factors may contribute to differences in the
   data?

6. How does your data for Scioto river water compare to values obtained by the City of Columbus for
   the influent to the Dublin Road water treatment plant? What factors may contribute to differences in
   the data?




                                                                                                    30
                                            LABORATORY #6

          COLORIMETRIC METHOD FOR THE DETERMINATION OF PHOSPHATE

Purpose

The objective of this laboratory is to determine the concentration of phosphorus in water using the
stannous chloride colorimetric method.

Background

Phosphate in natural waters. Phosphorus is an essential nutrient for algal growth, and when in excess it is
one of the leading causes of eutrophication. The primary sources of phosphorus in natural systems
include wastewater treatment facilities, runoff of fertilizer from agricultural operations, detergents and
some natural sources. Under summer growing conditions, it has been established that the critical level for
inorganic phosphorus is approximately 5 g/L.

Orthophosphates and polyphosphates are the most common forms of inorganic phosphorus found in
natural waters. Orthophosphates contain a single phosphorus molecule, and common orthophosphates
include trisodium phosphate (Na 3PO4), disodium phosphate (Na2HPO4), monosodium phosphate
(NaH2PO4), and diammonium phosphate ((NH4)2HPO4). Polyphosphates contain multiple phosphorus
molecules, and examples include sodium hexametaphosphate (Na3(PO3)6), sodium tripolyphosphate
(Na5P3O10), and tetrasodium pyrophosphate (Na4P2O7). Polyphosphates hydrolyze in natural waters to
the ortho form, typically in the time frame of several hours to days.

Several techniques are available for the determination of phosphorus in natural water samples, including
gravimetric, volumetric, and colorimetric methods. Gravimetric and volumetric methods are best when
the concentration of phosphorous is high. In most environmental engineering applications, this is not the
case, and colorimetric methods are preferred. The method detection level for phosphate by colorimetry is
approximately 0.1 mg/L as phosphorous, or 0.01mg/L if an extraction step is included, and thus it is the
most common technique used for the analysis of water and wastewater.

Colorimetry and Beers Law. Spectrophotometric methods are used to determine the concentration of
dissolved substances that absorb light in the ultraviolet (~180-400nm) or visible (~400-800nm)
wavelength range. These methods make use of the Beer-Lambert Law, i.e. the relationship of the amount
of light transmitted by a solution to the concentration of the light absorbing constituent, namely,

        Log(I0 / I)  A  bc                                                                            [1]

where I is the intensity of monochromatic light transmitted through the test solution, I 0 is the intensity of
light transmitted through the reference solution (or “blank”), A is the absorbance,  is the molar absorp-
tivity or extinction coefficient (a constant for a given solute and wavelength), b is the pathlength, and c is
the concentration of absorbing solute.

For solutions that obey Beer’s Law, light absorbance is directly proportional to the path length and the
concentration of the absorbing species. Alternatively, the intensity of the light transmitted through the
solution can be measured. The transmittance, T, of a solution is defined as I/ I 0 , and percent transmit-
tance as 100xT.




                                                                                                           31
In this set of experiments you will perform a spectrophotometric determination of orthophosphate in
solution. The instrument that will be used in this experiment is a visible wavelength spectrophotometer.
Light from a tungsten lamp is directed through a diffraction grating and a wavelength bandwidth of
~20nm is selected with a slit. The monochromatic beam is then passed through the sample, and light that
is not absorbed is received by a light-sensitive phototube. The analog or digital signal is then displayed
on the instrument panel. Variations at different wavelengths in the light source emission intensities,
response characteristics of the phototube, and light absorption by the cuvette, require that the intensity
control be adjusted to compensate for these effects.

The sensitivity of a spectrophotometric analysis depends on the degree of adherence to Beer’s Law, the
absorptivity of the species being measured relative to other compounds present, and the pathlength of the
sample cuvette. Deviations from Beer’s Law can occur for several reasons, particularly if the molar
absorptivity over the bandwidth of wavelengths varies significantly, if the sensitivity of the phototube is
exceeded, if the extent of equilibrium among species changes, or if one of the solutes fluoresces. Other
nonlinear effects are also possible. Spectrophotometric methods are most sensitive when the species of
interest has a much greater extinction coefficient than all other sample constituents. Furthermore, to
maximize the sensitivity of a measurement, wavelength selection should be determined from the absorp-
tion spectrum; the analyst must choose a wavelength near the region where maximum light absorption
occurs and where the absorptivity is reasonably constant over the 20nm bandwidth.

Stannous Chloride Colorimetric Method. In this laboratory we will use the stannous chloride colorimetric
method for the determination of orthophosphate in water samples. In this method, phosphates combine
with ammonium molybdate under acid conditions to form a molybdophosphate complex,

     PO43- + 12(NH4)2MoO4 + 24H+  (NH4)3PO412MoO3 + 21NH4+ + 12H2O

When large amounts of phosphate are present this reaction forms a yellow precipitate that can easily be
detected using a spectrophotometer. For cases in which the concentration of phosphate is low (<30mg/L),
which is common for most environmental applications, an additional step is needed. When stannous
chloride is added, it reduces the molybdophosphate complex forming a molybdenum blue complex,

    (NH4)3PO412MoO3 + Sn2+  (molybdenum blue) + Sn4+

The reduced blue complex is easily detected at 690nm using a standard UV-Vis spectrophotometer.
When polyphosphates are present, an additional acid hydrolysis step converts all polyphosphates to
orthophosphates. However, in most natural samples, orthophosphates represent greater than 95% of the
total phosphorus present.

Materials and Methods

In this laboratory, you will work in groups of 4-6 people. The following equipment and supplies are
required for completion of the laboratory.

       Glassware. (200mL volumetric flask [1 per analysis team], 500mL or 250 mL
        Erlenmeyer flasks [1 per standard], 10mL Mohr pipette, 2-3 clean spectrophotometric
        cuvettes, 100mL graduated cylinder (1), 250mL beaker (1-2 per analysis team), timer.

       Spectrophotometer. (capable of operating at a wavelength of 690nm)

The teaching assistant has prepared the following solutions.



                                                                                                        32
       Phenolphthalein Indicator. Prepared by the TA by dissolving 0.2g phenolphthalein into
        200mL deionized water and 200mL ethanol.

       Molybdate Reagent. Dissolve 12.5g of (NH4)6Mo7O244H2O in 100mL Deionized
        water. Cautiously add 140mL concentrated sulfuric acid (H2SO4) to 200mL Deionized
        water. Let the sulfuric acid cool (about 20 minutes) and then add the molybdate mixture
        to the sulfuric acid and dilute to 1L. Note: Always wear gloves, apron, and goggles
        when working with a strong acid-like sulfuric acid.

       Stannous Chloride Reagent. Dissolve 2.5g of fresh stannous chloride in 100mL
        glycerol (also known as glycerin). Heat on a hot plate (lowest setting) and stir with a
        stirring rod to enhance dissolution.

       Standard Phosphate Solution. Dissolve 219.5mg of potassium biphosphate (KH2PO4)
        and dilute to 1000mL (1mL=50.0 gPO43- - P).

Determine the amount of phosphate in tap water, Scioto river water, Olentangy river water, one blank,
and one duplicate using the stannous chloride colorimetric approach. Carry out the following procedures:

1. The spectrophotometers must be warmed up for at least 15 minutes prior to use, therefore make sure
   the spectrophotometer is turned “on”. Use the knob on the front-left side for the old instruments and
   the switch on the back for the new instruments.

2. Prepare calibration standards. Use a Mohr pipette to add the following volumes of the standard
   phosphate solution to 200 mL (or 250 mL) volumetric flasks. Fill to the mark with Milli-Q water.

              Standard 1:   0.0 mL (0 g/200 mL = 0.00 mg/L PO43- -P)
              Standard 2:   0.5 mL (25 g/200 mL = 0.125 mg/L PO43- -P)
              Standard 3:   1.0 mL (50 g/200 mL = 0.25 mg/L PO43- -P)
              Standard 4:   2.0 mL (100 g/200 mL = 0.5 mg/L PO43- -P)
              Standard 5:   4.0 mL (200 g/200 mL = 1.0 mg/L PO43- -P)

    Note. Remember to adjust phosphate concentrations if using a 250 mL flask (i.e., for Standard
    2 the correct 250 mL concentration is 0.100 mg/L PO43- -P).

3. Remove 100mL of each phosphate standard from the volumetric flask and place in an Erlenmeyer
   flask. Label these flasks as “Standard” with the g P, group name, and date. Note: depending upon
   your flask size, the g P will be half or 2/5 of your total addition.

4. Measure 100mL of each water sample to be analyzed and place in an Erlenmeyer flask.

5. To each flask, add 2 drops of phenolphthalein indicator. If a red color develops, add strong acid
   dropwise to discharge the color.

6. Add, with thorough mixing after each addition, 4.0mL of molybdate reagent and 10 drops of stannous
   chloride reagent.

7. After 10 minutes, but before 12 minutes, using the same interval for each sample, measure the color
   spectrophotometrically at 690nm. To zero the reading for the old instruments, use the left hand knob
   to set the transmittance to zero when no sample is in the instrument. Then place a sample of the



                                                                                                     33
    blank into the instrument and set the transmittance to 100% using the knob on the front-right side of
    the instrument. Change the mode to “Absorbance” and measure your sample. The new instruments
    are zeroed using the blank by pressing the “0 ABS” button. Always run the blank on a sample that
    includes the reagents.

Pre-Lab Questions

1. What is the difference between orthophosphate and polyphosphate and which type of phosphorus will
   the stannous chloride colorimetric method measure?

2. Show that if 219.5 mg of KH 2PO4 is added to 1 liter of water, the concentration of P in the solution is
   50 mg/L.

3. Given the following calibration data, determine the linear least squares best fit line to the data and the
   amount of phosphate in the unknown sample. If the sample size for the unknown was 100 mL, what
   is the concentration in mg/L?


              Sample                       Phosphate                  Absorbance
                                             (g)
              Standard 1                      25                           0.05
              Standard 2                      50                           0.11
              Standard 3                      100                          0.19
              Standard 4                      200                          0.42
              Standard 5                      400                          0.76
              Unknown                          ?                           0.35

Post-Lab Analysis

1. For your calibration standards, plot your absorbance values versus the amount of phosphate (g).
   Determine the linear least squares, best-fit line for the data and include this in your plot. This can be
   done using Excel and then taping the graph into your laboratory notebook.

2. What is a practical detection limit for phosphate based on the stannous chloride procedure you ran?

3. Using the fitted line (determined above), calculate the phosphate concentrations of the unknown
   samples. Based on these data, determine the concentration of phosphate (mg/L) in each sample as,


        PO  3
             4      
                  P 
                         μg PO3  P
                              4
                         mL sample




                                                                                                          34
4. Prepare a table summarizing your data as shown below:


      Sample                    Absorbance      Phosphate     Phosphate
                                                  (g)         (mg/L)
      Standard 1
      Standard 2
      Standard 3
      Standard 4
      Standard 5
      Blank
      Tap
      Scioto
      Olentangy
      (Replicate)

5. What is the precision of your analyses based on your duplicate samples?

6. How does your data for tap water compare to values obtained by the City of Columbus for the Dublin
   Road and Hap Creamen water treatment plants? What factors may contribute to differences in the
   data?

7. How does your data for Scioto river water compare to values obtained by the City of Columbus for
   the influent to the Dublin Road water treatment plant? What factors may contribute to differences in
   the data?




                                                                                                    35
                                           LABORATORY #7

                DISSOLVED OXYGEN AND BIOCHEMICAL OXYGEN DEMAND

Purpose

The objectives of this laboratory are to (1) measure the dissolved oxygen concentration of a water sample
using the azide modification of the Winkler procedure, and (2) calculate the BOD based on previously
determined dissolved oxygen measurements.

Background

Measurement of dissolved oxygen (DO) and biochemical oxygen demand (BOD) are two of the most
common tests to assess pollution in natural waters.

Dissolved Oxygen. Natural levels of dissolved oxygen in surface waters range from 7 mg/L to 14 mg/L,
depending on temperature, salt concentration, and the amount of biodegradable organic matter. When
organic pollution is present, for example due to a combined sewer overflow, microorganisms in the water
utilize the available oxygen to convert the organic material to cell mass and carbon dioxide. As a result,
the dissolved oxygen concentration can drop to levels significantly below 7 mg/L.

The two most common approaches for measuring dissolved oxygen in natural waters include the Winkler
method with Azide modification, and a membrane electrode method. In the Azide-modified Winkler
procedure, oxygen is fixed using manganous sulfate (MnSO44H2O), which is then used to oxidize iodide
(I-) to iodine (I2). The resulting I2, which will be proportional to the amount of O2 originally present, can
then be determined directly by titration with sodium thiosulfate.

The fixation of O2 occurs as,

        Mn2+ + 2OH- + ½ O2  MnO2 + H2O                                                                  [1]

This reaction is slow, and therefore, vigorous shaking for at least 20 seconds helps precipitate the white
manganese solid. The resulting precipitate then reacts with I-, in the presence of a strong acid, to produce
I2,

        MnO2 + 2I- + 4H+  I2 + 2H2O + Mn2+                                                              [2]

In this reaction, Mn4+ is reduced to Mn2+ while I- is oxidized to form I2. The iodine is usually complexed
with excess iodide to form a tri-iodide complex. The amount of I2 (or tri-iodide) in solution is then
determined by titration with sodium thiosulfate,

        2Na2S2O35H2O + I2  Na2S4O6 + 2NaI + 5H2O                                                       [3]

Thus, each mole of O2 present in the original sample will consume 4 moles of sodium thiosulfate.
Typically, starch is used as an indicator for the above reaction. Starch forms a dark blue color when
complexed with tri-iodide. When sodium thiosulfate has reacted with all the iodine, the starch complex
will be destroyed and the solution will turn clear.

Other compounds commonly found in water may also oxidize I- in the procedure above, most notably
nitrite (NO2-). Nitrite interference is easily overcome, however, by treating the sample, prior or during the



                                                                                                          36
addition of iodide, with sodium azide (NaN3). Under acid conditions, sodium azide reacts with any free
nitrite in solution producing N2 gas that leaves the aqueous phase,

       NaN3 + H+  HN3 + Na+                                                                                    [4a]

       HN3 + NO2- + H+  N2 + N2O + H2O                                                                         [4b]

Biochemical Oxygen Demand. The BOD test is one of the most common measures of organic matter in
wastewater and sewage-contaminated natural waters. In the BOD test, the amount of oxygen used in the
metabolism of biodegradable organics is termed the biochemical oxygen demand, or “BOD.” The
principal forms of biodegradable organic matter include proteins, carbohydrates, lipids, and fats.

The BOD of a water sample is determined by placing aliquots with appropriate dilution water in glass-
stoppered 300mL BOD bottles, incubating the bottles at a standard temperature (20°C), and measuring
the change in oxygen concentration with time. The concentration of dissolved oxygen in the BOD bottles
is determined using either a dissolved oxygen electrode or by performing the Winkler procedure.

The BOD, as a function of time, is assumed to follow a first-order rate model. Based on this model, the
BOD consumed or exerted in the BOD bottle at any time is equal to the difference between the BOD
existing at the initial time (BODu or Lo) and the BOD remaining at any time, (BODr or Lt).

        x  Lo  Lt  BOD u 1  e  kt                                                                         [5]

where x is BOD exerted at any time, t, BODu is the ultimate BOD (assumed equal to the oxygen
equivalent of organics at time zero, Lo), k is a first-order rate constant, and t is time.

                                            300
          Oxygen Equivalent BOD, L (mg/L)




                                                      Lo
                                            250

                                                           Lo - L t
                                            200                                         BOD Exerted


                                            150

                                                                       BODt
                                            100


                                            50                Lt                        L Remaining

                                             0
                                                  0                5        10          15            20   25
                                                                              Time (days)


Figure 4. Changes in biodegradable organics, measured in oxygen equivalents, as a
function of time.

Because BOD changes with respect to time, it is important to report at what time a BOD measurement
was made. The most common time interval for reporting BOD values is 5 days. For example, a number
of national standards for BOD are based on BOD5. The BOD is calculated as,


                                                                                                                  37
                D1  D2  f (B1  B2 )
        BOD                                                                                          [6]
                         P

where D1 and D2 are the dissolved oxygen concentrations of the mixtures in the BOD bottles before and
after incubation, respectively, B 1 and B2 are the dissolved oxygen concentrations in the dilution water
before and after incubation, respectively, f is the faction of dilution water in the mixture and P is the
fraction of sample in the mixture.

Materials and Methods

In this laboratory, you will work in groups of 4-6 people. The following equipment and supplies are
required for completion of the laboratory.

       Glassware. (500mL Erlenmeyer flask, 300mL BOD bottles, 10 or 15mL Mohr pipette,
        50mL burette)
       Manganous sulfate. Dissolve 480g of MnSO44H2O in deionized water and dilute to 1L.
       Alkali-iodide-azide. Dissolve 500g NaOH and 135g NaI in deionized water and dilute
        to 1L. Dissolve 10g NaN3 in 40mL deionized water and add it to the alkali-iodide
        solution.
       Sulfuric acid. Concentrated.
       Starch Indicator. Dissolve 2g of starch and 0.2g salicylic acid in 100mL of hot
        deionized water.
       Sodium thiosulfate. Dissolve 6.205g of Na2S2O35H2O in deionized water, add 0.4g
        NaOH and dilute to 1L. (Note: Sodium thiosulfate absorbs and loses water readily, so
        this solution should be standardized against potassium bi-iodate.)

Determine the amount of dissolved oxygen in tap water, Olentangy river water, and one duplicate using
the Winkler procedure. Carry out the following;

1. Obtain two samples of Olentangy river water in 300 mL BOD bottles according to standard methods,
   weather permitting. We will also measure the DO in the Olentangy river at this time using a
   dissolved oxygen meter. In the event of poor weather, samples and DO meter readings will be
   provided for you.

2. Fill cleaned and rinsed BOD bottles with each of your samples for analysis. The bottles should be
   filled close to the rim but not overflowing. Place the glass stopper in each bottle to remove any
   excess sample. Also, fill one BOD bottle with tap water to use as a blank.

3. Open each BOD bottle and add 1mL of manganous sulfate solution followed by 1mL of alkali-iodide-
   azide solution. Do this using a Mohr pipette. In order to prevent sample aeration due to the jet issuing
   from the pipette; hold the pipette tip under the surface of the sample. Thoroughly rinse the pipette
   interior and surface with deionized water after each use.

4. Re-stopper the BOD bottles carefully so as to exclude air bubbles and mix them by repeated inverting
   about 15 times. Note: Some of the original sample will be displaced when the stopper is reinserted.
   See below for correction.

5. Let the precipitate settle and repeat the mixing. Upon settling for the second time add 2mL of
   concentrated sulfuric acid. Again, prevent aeration of the bottle by letting the acid run down the




                                                                                                        38
    inside surface of the bottle neck. Re-stopper the bottle and mix several times to dissolve the
    precipitate.

6. Take enough liquid to represent 200mL of the original sample. Taking into account the 2mL of
   reagent-displaced sample, place approximately 201mL of the diluted sample into a 500mL
   Erlenmeyer flask, and titrate with 0.025N sodium thiosulfate until a pale straw color is obtained.

7. Add enough of the starch indicator to turn the solution dark blue (approximately 6 drops).

8. Titrate until the blue color disappears and the solution is completely clear. The color will return if the
   solution is let to stand, so report the first disappearance.

9. Calculate the dissolved oxygen as

        1 mg/L DO = 1 mL titrant

Pre-Lab Questions

1. Show that for the titration of a 200 mL sample, 1.0 mL of 0.025 M Na2S2O3 is equivalent to 1.0 mg/L
   dissolved oxygen.

2. Calculate the dissolved oxygen of a sample if 200 mL requires 8.4 mL 0.025 Na2S2O3 to reach the
   titrimetric endpoint.

3. What is the 7-day BOD if the initial DO concentration of the mixture is 10.2 mg/L, the DO
   concentration of the mixture after 7 days is 3.4 mg/L, the initial DO concentration of the dilution
   water is 9.8 mg/L, the DO concentration of the dilution water after 7 days is 9.2 mg/L, the fraction of
   dilution water in the mixture is 0.1 and the fraction of sample in the mixture is 0.9.

4. Derive the equation for BOD as a function of time (Eq. 5).

5. Assuming a reaction rate constant of 0.23 (day-1), determine the 5-day BOD if the 7-day BOD is 52
   mg/L.

Post-Lab Analysis

1. Calculate the dissolved oxygen concentrations of your samples. Prepare a table summarizing your
   data as shown below:


                                       (Meter)             Volume           (Winkler)
                                  Dissolved Oxygen         Titrant      Dissolved Oxygen
               Sample                  (mg/L)               (mL)             (mg/L)
               Tap
               Olentangy-1
               Olentangy -2

2. What is the precision of your analyses based on your duplicate samples?

3. How do your dissolved oxygen measurements based on the Winkler method compare to data
   determined using a dissolved oxygen meter? What factors may account for any differences observed?



                                                                                                          39
                                           LABORATORY #8

                                      COLORIMETRY OF IRON

Purpose

The objective of this laboratory is to determine the amount of total dissolved iron in a water sample using
a colorimetric approach.

Background and Theory

Iron in Natural Waters. Iron is a common element in soil and is ubiquitous in natural waters. Common
iron minerals include iron hydroxides (e.g., hematite, goethite), iron sulfides (e.g., pyrite), and iron
carbonates (e.g., siderite). In lakes, rivers, and groundwater iron may be present in either the ferrous
(Fe2+) or ferric (Fe3+) form depending on the oxidation-reduction conditions. In well-oxygenated surface
waters, the ferric form of iron will predominate, whereas in surface waters impacted by high levels of
organic matter, or low dissolved oxygen ground waters, ferrous iron may also be important. The ferrous
form is more soluble than the ferric form, and therefore, anaerobic conditions generally result in increased
levels of total soluble iron.

Although moderate levels of iron in drinking water do not have significant health effects, the United
States Environmental Protection Agency has promulgated a secondary maximum contaminant level for
iron of 0.3 mg/L. Moderate levels of iron create drinking water taste problems (bittersweet, astringent
taste) and can stain sinks, bathtubs and other plumbing fixtures. Soluble iron precipitates as iron
hydroxide (Fe(OH)3) producing a characteristic orange or red color. The most common approaches for
determining the amount of iron in natural waters include atomic absorbance spectroscopy, inductively-
coupled plasma (ICP) atomic emission spectroscopy, and the phenanthroline colorimetric method.

Colorimetry and Beers Law. See Laboratory #6, p. 30.

Phenanthroline Colorimetric Method for Iron. Dissolved iron does not absorb in the ultraviolet or visible
range and cannot be determined directly by colorimetric methods. Therefore, in this method, soluble
ferrous iron is reacted with 1,10-phenanthroline to produce an orange-colored complex. It is important to
note that 1,10-phenanthroline only reacts with ferrous iron. Therefore, if ferric or total dissolved iron is
of interest, the iron in the samples must first be reduced from the ferric to the ferrous form.

The following approach can be used to determine total soluble iron in a water sample (ferric+ferrous). In
the first step hydroxylamine is added which reduces all ferric iron to ferrous iron,

4Fe3+ + 2NH2OH  4Fe2+ +N2O +4H+                                                                        [1]

Next, 1,10-phenanthroline is added which reacts with the ferrous iron and forms an orange-red complex,

Fe2+ + 3(1,10-phenanthroline)  Fe-(1,10-phenanthroline)3                                               [2]

The amount of iron in the original sample is then directly proportional to the intensity of the iron-
phenanthroline complex determined spectrophotometrically.




                                                                                                         40
Materials and Methods

In this laboratory, you will work in groups of 4-6 people. The following solutions have been prepared
prior to laboratory.

       Glassware. (100mL volumetric flask (5 or 10), 50mL volumetric pipette, 1L volumetric
        flask, 10mL volumetric pipette, 5mL volumetric pipette, 1mL Mohr pipette, 10mL Mohr
        pipette, clean cuvettes (5))

       Concentrated hydrochloric acid. The acid must contain less than 0.00005% iron.

       Ammonium acetate buffer solution. Dissolve 250g ammonium acetate [NH4CH3COO]
        in 150mL deionized water. Add 700mL glacial acetic acid and bring up to 1 L with
        deionized water.

       Hydroxylamine solution. Dissolve 10g NH2OH HC1 in 100mL of deionized water.

       0.1 N Potassium Permanganate solution. Dissolve 3.2g of potassium permanganate
        [KMnO4] in 1L of deionized water.

       Phenanthroline solution. Dissolve 1.0g of 1, 10 phenanthroline monohydrate
        [C12H8N2(H2O)] in 1L of deionized water that contains 20 drops of concentrated hydro-
        chloric acid.

Note: The stoichiometric ratio for the complex is 3 moles of phenanthroline per mole of Fe 2+, or 10.65g
phenanthroline per g of Fe 2+. This implies that 1mL (1 mg phenanthroline) of this reagent is sufficient for
about 100  g of Fe. However, in order to get rapid, complete complexation, excess phenanthroline is
required and in practice 1mL of this solution suffices for no more than 2.5  g of Fe2+.

       Stock Iron Solution. Slowly add 20mL of concentrated sulfuric acid to 50mL of deionized
        water and dissolve 1.404g of ferrous ammonium sulfate [Fe(NH4)2(SO4)2(H2O)6] in the acid/water
        mixture. Add 0.1 N potassium permanganate dropwise until a faint pink color persists. Dilute to
        1L with deionized water. Note, 1.00mL=200  g Fe.

       Standard Iron Solution: Prepare this solution daily. Using a volumetric pipette, pipette 50mL
        of the stock iron solution into a 1L volumetric flask and dilute to 1L with deionized water. Note,
        1.00mL=10  g Fe.

Determine the amount of total dissolved iron in samples of Olentangy river water, Scioto river water, tap
water, one blank, one check standard, and one duplicate. Carry out the following procedures:

1. Review spectrometer warm-up procedures in Laboratory #6.

2. Prepare the standard iron solution. Using a volumetric pipette, pipette 50mL of the stock iron solution
   into a 1L volumetric flask and dilute to 1L with deionized water. Note, 1.00mL=10  g Fe.

3. Prepare the following samples. Set up the first 5 100mL volumetric flasks, and use a volumetric
   pipette to add the following amounts of standard iron solution or sample to the flasks (note: if flasks




                                                                                                         41
    are limited than you will need to run your samples in two batches. Do flasks 1-5 first, then run flasks
    6-10):

                Flask 1: none [the “blank”=0.00 g/L]
                Flask 2: 5.0mL standard iron [50  g/100mL = 0.5mg/L Fe]
                Flask 3: 10.0mL standard iron [100  g/100mL = 1.0mg/L Fe]
                Flask 4: 20.0mL standard iron [200  g/100mL = 2.0mg/L Fe]
                Flask 5: 30.0mL standard iron [300  g/100mL = 3.0mg/L Fe]
                Flask 6: 50.0 mL Olentangy river water
                Flask 7: 50.0 mL Olentangy river water
                Flask 8: 50.0 mL Scioto river water
                Flask 9: 50.0 mL Tap water
                Flask 10: Olentangy river water plus 2.0 mg/L Fe

4. Using measuring pipettes, add 2mL of concentrated hydrochloride acid, 1mL of the hydroxylamine
   solution, 10.0mL of the ammonium acetate solution, and 10mL of the phenanthroline solution to each
   flask. (Note: This quantity of phenanthroline suffices for 1mg of Fe, which is equal to 100mL of
   standard iron solution.)

5. Add enough deionized water to bring the volume in each flask to 100mL. Mix the flasks thoroughly
   and let them stand for 20 minutes to develop the reddish-orange color. (Note: The blank may be
   colorless.)

6. If needed, review spectrometer operating procedures in Laboratory #6. Set the wavelength of the
   spectrophotometer to 510nm and after 20 minutes, using the same time interval for each sample,
   measure the absorbance. Always start analyses with the “blank”.

Pre-Lab Questions

1. What is the difference between ferric and ferrous iron? Describe how each can be measured using the
   phenanthroline colorimetric method.

2. Given the following calibration data, determine the least squares best fit line to the data and determine
   the amount of iron in the unknown sample. If the volume of the unknown sample used in the analysis
   was 50 mL, what is the concentration?


              Sample                          Iron                    Absorbance
                                              (g)
              Standard 1                       50                         0.12
              Standard 2                      100                         0.21
              Standard 3                      200                         0.39
              Standard 4                      300                         0.62
              Standard 5                      400                         0.76
              Unknown                           ?                         0.35




                                                                                                         42
Post-Lab Analysis

1. For your calibration standards, plot your absorbance values on the x-axis versus mass iron (g) on the
   y-axis. Determine the linear least squares, best-fit line for the data and include this in your plot. This
   can be done using Excel and then taping the graph into your laboratory notebook.

2. The reported MDL for iron based on the phenanthroline procedure is 50 g/L. Does this correspond
   to your measurements? What is the practical detection limit based on the reported MDL?

3. Using the fitted line (determined above), calculate the amount of iron, and the iron concentration of
   the unknown samples. The concentration of iron (in mg/L) is determined from the mass of iron as,

                             μg Fe
                 Fe  
                           mL sample

4. Prepare a table summarizing your data as shown below:


               Sample            Absorbance           Iron           Iron
                                                      (g)          (mg/L)
               Blank
               Tap
               Scioto
               Spike
               Olentangy-1
               Olentangy -2



5. Is your blank within appropriate limits? If not, what errors may have been introduced into your
   analysis to account for this?What is the precision of your analyses based on your duplicate samples?

6. Calculate the percent recovery of your spike sample. Is it acceptable? If not, what errors may have
   been introduced into your analysis to account for this?

7. How does your data for tap water compare to values obtained by the City of Columbus for the Dublin
   Road and Hap Creamen water treatment plants? What factors may contribute to differences in the
   data.

8. How does your data for Scioto river water compare to values obtained by the City of Columbus for
   the influent to the Dublin Road water treatment plant? What factors may contribute to differences in
   the data?




                                                                                                          43
Acknowledgment

A number of individuals have contributed to this laboratory manual. Many of the experiments have been
adapted from previous laboratory courses taught by Professors H.W. Walker, R.M. Sykes, and A.J. Rubin
at The Ohio State University. The experiments for determining Hardness and Solids were adapted from
Standard Methods for the Examination of Water and Wastewater. Suggestions made by undergraduates
and laboratory assistants have been especially useful in improving this manual.




                                                                                                  44

				
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